simple mixtures chapter 5 outline hw: sect. title and...

SIMPLE MIXTURES

Chapter 5 Outline HW: Questions are below. Solutions are in separate file on the course web site. Sect. Title and Comments Required?

1. Partial Molar Quantities YES 2. The Thermodynamics of Mixing YES 3. The Chemical Potentials of Liquids YES 4. Liquid Mixtures MOST

We will not cover Sect. 4.b on Excess Functions and Regular Solutions.

5. Colligative Properties YES 6. Vapor Pressure Diagrams NO 7. Temperature-Compositon Diagrams NO 8. Liquid-Liquid Phase Diagrams NO 9. Liquid-Solid Phase Diagrams NO 10. The Solvent Activity A LITTLE I’ll just discuss real solutions briefly, introducing activities and activity coefficients. We won’t perform any calculations on real solutions. 11. The Solute Activity A LITTLE I’ll just discuss real solutions briefly, introducing activities and activity coefficients. We won’t perform any calculations on real solutions. 12. The Activities of Regular Solutions NO 13.. The Activities of Ions in Solution NO

Chapter 5 Homework Questions

5.1 At 25 oC, the density of a 50% by mass of an Ethanol-Water solution is 0.914 g/cm3. Given that the Partial Molar Volume of water in the solution is 17.4 cm3/mol, calculate the Partial Molar Volume of Ethanol in the solution.

Note: M(Ethanol) = 46. g/mol , M(Water) = 18. g/mol. 5.2 The vapor pressure of pure benzene (C6H6, M=78) is 53.3 kPa at 60 oC. When 19. grams

of an involatile organic compound is dissolved in 500 g of Benzene, the vapor pressure drops to 51.5 kPa.

Calculate the Molar Mass of the organic compound. 5.3 The freezing point of pure CCl4(liq) is -22.9 oC and the Freezing Point Depression

constant is 30 oC. When 100 grams of an unknown organic compound is added to 750 grams of CCl4(l), the freezing point of the mixture is -33.4 oC

Calculate the Molar Mass of the organic compound. 5.4 The boiling point of pure benzene is 80.1 oC and the Boiling Point Elevation constant is 2.13 oC/m. When a sample of napthalene (C10H8) is dissolved in 600. grams of

Benzene, the boiling point boiling point of the mixture is 81.3 oC. How many grams of napthalene were dissolved in the benzene. 5.5 When 0.15 grams of an unknown compound is dissolved in 100 mL of aqueous solution,

the measured osmotic pressure of the solution is 0.65 kPa at 25 oC. Calculate the molar mass of the unknown compound.

5.6 Consider two containers separated by a partition. Container A is of volume 5 L, and

contains N2(g) at 2.0 atm and 30 oC. Container B is of volume 10 L, and contains H2(g) at 2.0 atm and 30 oC.

Calculate the Entropy of mixing and the Gibbs Energy of mixing when the partition between the two partitions is removed.

5.7 Air is a mixture of primarily 3 gases with composition: xN2 = 0.78 , xO2 = 0.21 , xAr = 0.01

Calculate the Entropy of mixing when 5 moles of air of the above composition is prepared from the above pure gases. 5.8 The freezing point of 1-butanol is 25.8 oC and its depression constant is 8.2 oC/m. When 4.0 grams of acetonitrile (CH3CN, M = 41) is dissolved in 650 grams of 1-butanol, the freezing point of the mixture is 21.5 oC.

Calculate the activity coefficient of acetonitrile in 1-butanol.

1

Chapter 5

Simple Mixtures

Chapter 7 : Slide 1

Partial Molar Quantities

Partial Molar Volume

Imagine adding 1 mole (18 g) of H2O(l) to a very large volume of water.The volume would increase by 18 cm3, and we would say that

and would call this the Partial Molar Volume of water.

If, instead, we added 1 mole of water to a very large volume of ethanol,the volume would increase by only 14 cm3 because of packing effects.Under these conditions, the Partial Molar Volume of water is 14 cm3/mol.

Thus, the volume increase depends upon the nature of the solution; i.e.the number of moles of the various components.

In general, the Partial Molar Volume, VJ , of a substance J is defined as:

n' indicates that the numbers of moles of all componentsexcept J are held constant.

Chapter 7 : Slide 2

2

Dependence of VH2O and VEtOH on composition in a binarywater/ethanol solution.

Chapter 7 : Slide 3

For a binary mixture of A and B, the incremental change in volume isgiven by:

Integrate this expression under conditions of constant composition so that VA and VB are constant, to find that the total volume of the mixture, V, is given by:

This is may be generalized to an N component system:

and

n' indicates that all nj ni are held constant

Chapter 7 : Slide 4

3

Example: Exer. 5.2(b)

At 20 oC, the density of a 20% by mass solution of Ethanol and Water is0.97 g/cm3.

Given that the Partial Molar Volume of ethanol in this solution is 52.2 cm3/mol, what is the Partial Molar Volume of water?

M(EtOH) = 46 g/molM(H2O) = 18 g/mol

Hint: Assume 1000 cm3 of solution.

m(EtOH) = 194 g n(EtOH) = 4.22 molm(H2O) = 776 g n(H2O) = 43.1 mol

V(H2O) = 18.1 cm3/mol

Chapter 7 : Slide 5

Partial Molar Gibbs Energy

We introduced the Partial Molar Gibbs Energy, which is also calledthe Chemical Potential (), in Chapter 3.

where: n' indicates that all nj ni are held constant

If Pressure and Temperature are held constant, then one has:

Similar to Partial Molar Volumes, this expression can be integrated underconditions of constant composition (so that the i's remain constant) to yield:

Chapter 7 : Slide 6

4

The Gibbs-Duhem Equation

From above, we have:

if T, p are constant

We are going to develop an expression relating changes in the chemicalpotentials (di) of various components in a mixture.

We also have:

Taking the differential of the latter expression gives:

The two equations for dG must be equal:

which finally gives: Gibbs-Duhem Equation

Chapter 7 : Slide 7Note: You are not responsible for the derivation..

Chapter 7 : Slide 8

which finally gives: Gibbs-Duhem Equation**

Note: This is just a special case of the Gibbs-Duhem equation, which can bederived for any Partial Molar Quantities (e.g. the Partial Molar Volume, Vi)

The significance of the Gibbs-Duhem Equation is that the Chemical Potentialof one component of a mixture cannot change independently of the ChemicalPotentials of the other components.

For example, in a binary mixture (components A and B), one has:

This equation shows that if the chemical potential of one component increases,that of the other component must decrease; e.g.

The Gibbs-Duhem Equation for Chemical Potential and other quantitieshas important applications throughout thermodynamics.

Chapter 7 : Slide 8

5

The Thermodynamics of Mixing

The Chemical Potential of a Gas

In Chapter 3, we developed the following equation for the pressure dependenceof the Molar Gibbs Energy of a gas:

or, remembering that Gm

If we take the standard reference state as po = 1 bar, with (po) = o , thenthe chemical potential at any other pressure, p, is:

It is common (and convenient) to leave off the po, giving:

Chapter 7 : Slide 9

The Gibbs Energy of Mixing (mixG) of Perfect Gases

Consider two containers at the same temperature and pressure, each with a pure gas, A (nA moles) and B (nB moles)

We would like to consider the Gibbs Energy change when wemix these gases. The mixture is in a new container, with V = VA + VB , andtotal pressure, p = pA + pB .

The initial Gibbs Energy of the two gases in their separate containers is:

After mixing, the partial pressures of the two gases are pA and pB , withpA + pB = p. The final Gibbs Energy of the mixture is:

Chapter 7 : Slide 10

6

Initial:

Final:

Therefore, the Gibbs Energy change of mixing is given by:

With simplification:

We can simplify this further by using:

and Dalton's Law:


We can simplify this further by using:

and Dalton's Law:

xA mixG/nRT

0.1 -0.330.2 -0.500.3 -0.610.4 -0.670.5 -0.69 0.6 -0.670.7 -0.610.8 -0.500.9 -0.33


7

The Entropy of Mixing (mixS) of Perfect Gases

Therefore:

xA mix/nR

0.1 0.330.2 0.500.3 0.610.4 0.670.5 0.69 0.6 0.670.7 0.610.8 0.500.9 0.33


The Enthalpy of Mixing (mixH) of Perfect Gases

negative for all xA , xB

positive for all xA , xB

It is not at all surprising the the Enthalpy of Mixing is 0,because there are no attractive or repulsive forces in Perfect Gases.

Thus, we see that the spontaneous mixing of two gases is anentropically driven process.


8

+

A B

UAA UBB UAB

AB

Ideal Mixture: UAA UBB UAB

Nearly Ideal Mixtures: Benzene-TolueneHexane-Octane

Non-Ideal Mixtures: Chloroform-AcetoneEthanol-Water

The Chemical Potential of Liquids

Ideal Solutions: Raoult's Law

Chapter 7 : Slide 15Chapter 7 : Slide 15

pA* pB

* ptot = pA + pB

+

pA = xApA* xA = Mole Fraction of A in mixture

pB = xBpB* xB = Mole Fraction of B in mixture

ptot = xApA* + xBpB

*

Raoult’s Law

pA* is the vapor pressure above pure liquid A

pB* is the vapor pressure above pure liquid B

Chapter 7 : Slide 16Chapter 7 : Slide 16

9


The Chemical Potential in Ideal Solutions

We will use a superscript asterisk (*) to indicate a pure liquid. Thusthe chemical potential would be denoted as A

* or A*(l)

Because a pure liquid must be in equilibrium with its vapor, thechemical potential of the liquid is given by:

Ao is the chemical potential of the gas at po = 1 bar

pA* is the vapor pressure of the pure liquid

If another substance (solute) is present in the liquid, the chemical potentialof the liquid is changed to A and its vapor pressure is changed to pA

pure liquid

liquid in solution

pA is the vapor pressure of A in the solution


pure liquid

liquid in solution

We can combine the equations to eliminate Ao (subtract first from second):

or:

using Raoult's Law

Therefore, an ideal solution is one in which the chemical potential ofeach component is given by:

10


xB0 1

Pre

ssur

e

xA1 0

ptot = xApA* + xBpB

*

ptot

pB*

pB

pA*

pA

The Component Vapor Pressures in an Ideal Solution


A Nearly Ideal Solution

11


Positive Deviations from Raoult’s Law


Negative Deviations from Raoult’s Law

12


Ideal Dilute Solutions

Consider a solution with solvent, A, and solute, B.If it is an ideal solution, the vapor pressures of solvent and solute are given by:

and

It is found in many real dilute solutions (xB << xA) that, while the solvent vapor pressure is given by the expression above,the solute vapor pressure is still proportional to xB , but the proportionalityconstant is not pB

* , but an empirical constant.

and

This is called Henry's Law, and the Henry's Law constant, KB , is a functionof the nature of the solute, B.


The Properties of Solutions

Thermodynamics of Mixing of Ideal Solutions

If we have nA moles of A and nB moles of B, then the initial Gibbs Energy of the pure liquids before mixing is:

After mixing, the final Gibbs Energy of the solution is:

It is straightforward to show that the Gibbs Energy of Mixing is:

This is identical to the expression for the Gibbs Energy of Mixing ofPerfect Gases.

13


This is identical to the expression for the Gibbs Energy of Mixing ofPerfect Gases.

Similarly, it can be shown that the expressions formixS and mixH are the same as for Perfect Gas mixtures.

Excess Functions

Often, the mixing properties of Real solutions are discussed in terms ofexcess functions; e.g.

This is discussed in somewhat more detail in the text (Sect. 5.4b), but we will not go into excess functions any further.

Chapter 6: Slide 26

Colligative Properties

Colligative properties of a solution are properties which depend only uponthe number of moles of a solute, and not upon its specific nature.

• Vapor Pressure LoweringWhen an involatile (e.g. solid) solute is dissolved in a solvent,the vapor pressure drops below that of the pure solvent.

• Boiling Point ElevationWhen an involatile (e.g. solid) solute is dissolved in a solvent, the boiling point of the solution is higher than that of the pure solvent.

• Freezing Point DepressionWhen a solute (usually a solid) is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solvent.

• Osmotic PressureSolvent tends to flow from a pure solvent chamber into a solutionchamber until a sufficient pressure (the “Osmotic Pressure”) hasdeveloped to stop further flow.

14


Colligative Properties

All four colligative properties are independent of the nature of the solute (B), but stem from the fact that the solute will lower the chemical potential of the solvent (A):

Freezing PointDepression

Boiling PointElevation

pA* pB

* ptot = pA + pB

+

pA = xApA* xA = Mole Fraction of A in mixture

pB = xBpB* xB = Mole Fraction of B in mixture

ptot = xApA* + xBpB

*Raoult’s Law:

Vapor Pressure Lowering

In the special (but important) case that the solute, B, is non-volatile, it'spure vapor pressure will be: pB* 0

In this case, Raoult's Law reduces to: ptot = pA = xApA*


For xA < 1, pA < pA*

15


Application of Raoult's Law: Determining the Molar Mass of a compound.

Example: The vapor pressure of 2-propanol (M=60) is 50.0 Torr at 27 oC,but fell to 49.6 torr when 8.7 g of an involatile organic compound was dissolvedin 250 g of 2-propanol.

Calculate the Molar Mass of the unknown compound.

xprop = 0.992

nX = 0.034 mol

MX = 260 g/mol

Pure Solvent

p = pA* = 1 atm.

T = Tbo

Solution

P = xApA* < 1 atm.

T = Tbo

Will no longer boil at Tbo. Must raise

temperature until P rises back to 1 atm.

Boiling Point Elevation

Tb = Tb - Tbo = KbmB

It can be shown (numerous texts) that the change in the boiling point, Tb , can be expressed as a function of the solute molality, mB .

16

Pure Solvent

Rate(“pop off”) = Rate(“condense”)T = Tf

o

Solution

Will no longer freeze at Tfo. Must lower

temperature until the two rates are equal.

Rate(“pop off”) > Rate(“condense”)T = Tf

o

Tf = Tfo - Tf = KfmB

It has been shown in various texts that the change in the freezing point, Tf , can be expressed as a function of the solute molality, mB .

Freezing Point Depression


Tb = Tb - Tbo = KbmB

Tf = Tfo - Tf = KfmB

It can be shown that the Boiling Point Elevation Constant (aka the Ebullioscopicconstant) and the Freezing Point Depression Constant (Cryscopic constant) aregiven by:

and

MA is the solvent Molar MassTb

o and Tfo are the pure solvent boiling and freezing temperatures

vapHo and fusHo are the solvent enthalpies of vaporization and fusion

Note that Kb and Kf depend solely upon solvent properties, and independent of the particular solute.

Although Kf and Kb can be calculated from the above formulae, the tabulatedvalues in the literature are the empirical values, taken from experimental determination using solutes of known Molar Mass.

17


Solvent Kf Kb

H2O 1.86 K/m 0.51 K/m

Benzene 5.12 2.13

Phenol 7.27 3.04

CCl4 30 4.95

Note that Kf > Kb. This is because fusHo << vapHo

and

Generally, the tabulated constants, Kf and Kb , are the empirical valuesdetermined from measurements with solutes of known Molar Mass.


Applications of freezing point depression and boiling pointelevation measurements.

1. Determination of solute Molar Mass

2. Determination of fractional dissociation and equilibrium constants

3. Determination of solute "activity coefficients" (coming up)

mX = 0.112 m = 0.112 mol/kg

kg Nap = 0.25 kg

nX = 0.028 mol

MX = 180 g/mol

Example: The addition of 5.0 g of an unknown compound to 250 g ofNapthalene lowered the freezing point of the solvent by 0.78 oC


Kf (Nap) = 6.94 K/m

18

Osmotic Pressure

Wine bladder

LakePokeytenoqua

H2O


Chapter 6: Slide 36

SemipermeableMembrane


H2O

= dgh= Osmotic Pressure

Static Osmometry:Allow the two chambersto equilibrate.


19

Dynamic Osmometry: Apply pressure, , to prevent movementof solvent between chambers.

P P +


H2O


It can beshown that: V = nBRT

= (nB/V) RT

= Osmotic Pressure

V = Volume of solution

nB = Moles of Solute

R = Gas Constant

T = Temperature (in K)

Alternate form:

= [B]RT

[B] = Solute Molarity


20


A Sensitivity Comparison

Let's say that a sugar solution is prepared by dissolving 1 g of sucrose(M = 342) in 1 L (1 kg) of water.

The Molarity (molality) of the solution is:

From Kf = 1.86 K/m, one has Tf = -0.005 oC

From Kb = 0.51 K/m, one has Tb = 100.0015 oC

Let's calculate (25 oC)

In contrast to freezing point depression, and boiling point elevation, whichare too small to be measured experimentally, the osmotic pressure can be measured very accurately.


Application: Determination of Molar Mass from Osmotic Pressure

We will show a straightforward method for determining Molar Masses frommeasured values of . The method is very similar to that used to calculateMolar Mass from Freezing Point Depression or Boiling Point Elevation.

Afterwards, I will comment on the method presented in the text, whichis useful if one wishes very accurate determinations, but is not straightforward.

Example: When 2.0 g of Hemoglobin (Hb) is dissolved in 100 mLof solution, the osmotic pressure is 5.72 torr at 25 oC.

Calculate the Molar Mass of Hemoglobin. 1 kPa = 7.50 torrR = 8.31 kPa-L/mol-K

Procedure:

1. Calculate the Molarity of the unknown, [X] from = [X]RT

2. Determine the number of moles, nX , using [X] and the solution volume

3. Determine the Molar Mass from: MX =massX/nX

21


Example: When 2.0 g of Hemoglobin (Hb) is dissolved in 100 mLof solution, the osmotic pressure is 5.72 torr at 25 oC.

Calculate the Molar Mass of Hemoglobin. 1 kPa = 7.50 torrR = 8.31 kPa-L/mol-KProcedure:

1. Calculate the Molarity of the unknown, [X] from = [X]RT

2. Determine the number of moles, nX , using [X] and the solution volume

3. Determine the Molar Mass from: MX =massX/nX

Note: The Boiling Point elevation and Freezing Point depression of the abovesolution would be immeasurably small. Therefore, these techniques are uselessto determine Molar Masses of polymers.


Homework: When 3.0 g of an unknown compound is placed in 250 mL ofsolution at 30 oC, the osmotic pressure is 62.4 kPa.


R = 8.31 kPa-L/mol-K

[X] = 2.48x10-2 mol/L

nX = 6.20x10-3 mol

MX = 480 g/mol

In contrast, the freezing point depression (in aqueous solution) would be~0.05 oC, which could only be measured to approximately 10-20% accuracy.

Note: The osmotic pressure, = 62.4 kPa = 470 mm Hg 0.50 meters,can be measured to 0.1 % accuracy very easily.

Note, once again, that Osmotic Pressure measurements are much more sensitive than the other colligative properties.

22


FYI: Very Accurate Determination of Molar Masses from

Actually, the formula, = [B]RT, is only accurate in dilute solution.

In more concentrated solutions, one has:

It is shown in the text (Example 5.4) that, writing = gh (hydrostaticpressure), and c = mass/Volume, one can derive the equation:

Then, an accurate value for MX can be obtained from the intercept of a plotof h/c vs. c


Real Solutions: Activities

Ideal Solutions

Earlier in the chapter, we showed that the solvent (A) and solute (B) chemical potentials are given by:

Solvent:

Solute:

Real Solutions

In solutions which are not ideal, the relatively simple formulae for the solventand solute chemical potentials can be retained by replacing the mole fractionsby activities (ai):

Solvent:

Solute:

23


Real Solutions: Activities

Real Solutions

In solutions which are not ideal, the relatively simple formulae for the solventand solute chemical potentials can be retained by replacing the mole fractionsby activities (ai):

Solvent:

Solute:

The activities can be related to the mole fractions by:aA = A xA and aB = B xB.

A and B are the activity coefficients, and can be calculated fromthermodynamic measurements. Their deviation from i =1 characterizesthe extent of non-ideality in a system.


Solute Activities in terms of molalities

In a dilute solution, it is more common to characterize solutes by theirmolality, rather than mole fraction in solution.

It is straightforward to show that, in a dilute solution, the molality, mB,and mole fraction, xB , are directly proportional:

(nB << nA in dilute solution)

If we assume an amount of solution containing 1 kg = 1000 g of solvent,then: nB = mB and nA = 1000/MA (solvent Molar Mass).

24


The solute chemical potential of an ideal solution, written in terms of molalityrather than mole fraction, is:

B+ is different from the reference state, B

o ,when using mole fractions.

For a non-ideal solute, we can then replace mB by the activity, aB= BmB


Activity Coefficients from Colligative Property Measurements

Earlier we discussed that the freezing point depression and boiling point elevationof a solution (assumed ideal) are given by:

If the solution is non-ideal, the solute molality, mB , is replaced by itsactivity, aB

Thus, the activity coefficient, B , can be determined by a comparison ofthe measured colligative property with the value expected for an ideal solution:

25


Example: When 25 g of Napthalene (C10H8 , M = 128) is placed in 500 g of benzene (Kf = 5.12 oC/m. Tf

o = 5.5 oC ), the freezing point of thesolution is 3.9 oC.

What is the activity coefficient of napthalene in this solution?

simple mixtures chapter 5 outline hw: sect. title and...

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